How Atoms Work

Atoms are in your body, the chair you are sitting in, your desk and even in the air. See more nuclear power pictures.

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­It has been said that during the 20th century, man harnessed the power of the atom. We made atomic bombs and generated electricity by nuclear power. We even split the atom into smaller pieces called subatomic particles.

But what exactly is an atom? What is it made of? What does it look like? The pursuit of the structure of the atom has married many areas of chemistry and physics in perhaps one of the greatest contributions of modern science.In this article, we will follow this fascinating story of how discoveries in various fields of science resulted in our modern view of the atom. We will look at the consequences of knowing the atom's structure and how this structure will lead to new technologies.

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What Is an Atom? The Legacy of Ancient Times Through the 19th Century

The modern view of an atom has come from many fields of chemistry and physics. The idea of an atom came from ancient Greek science/philosophy and from the results of 18th and 19th century chemistry:

Repeat step 3 until you are down to a pile containing only one paper clip. That one paper clip still does the job of a paper clip (i.e., hold loose papers together).

Now, take a pair of scissors and cut that one paper clip in half. Can half of the paper clip do the same job as the single paper clip?

If you do the same thing with any element, you will reach an indivisible part that has the same properties of the element, like the single paper clip. This indivisible part is called an atom.

The idea of the atom was first devised by Democritus in 530 B.C. In 1808, an English school teacher and scientist named John Dalton proposed the modern atomic theory. Modern atomic theory simply states the following:

Every element is made of atoms - piles of paper clips.

All atoms of any element are the same - all the paper clips in the pile are the same size and color.

Atoms of different elements are different (size, properties) - like different sizes and colors of paper clips.

Atoms of different elements can combine to form compounds - you can link different sizes and colors of paper clips together to make new structures.

In chemical reactions, atoms are not made, destroyed, or changed - no new paper clips appear, no paper clips get lost and no paper clips change from one size/color to another.

In any compound, the numbers and kinds of atoms remain the same - the total number and types of paper clips that you start with are the same as when you finish.

Dalton's atomic theory formed the groundwork of chemistry at that time. Dalton envisioned atoms as tiny spheres with hooks on them. With these hooks, one atom could combine with another in definite proportions. But some elements could combine to make different compounds (e.g., hydrogen + oxygen could make water or hydrogen peroxide). So, he could not say anything about the numbers of each atom in the molecules of specific substances. Did water have one oxygen with one hydrogen or one oxygen with two hydrogens? This point was resolved when chemists figured out how to weigh atoms.

Important Terms

atom - smallest piece of an element that keeps its chemical properties

compound - substance that can be broken into elements by chemical reactions

molecule - smallest piece of a compound that keeps its chemical properties (made of two or more atoms)

neutron - particle in the nucleus of an atom with no charge (mass = 1.675 x 10-24 grams)

nucleus - dense, central core of an atom (made of protons and neutrons)

proton - particle in the nucleus of an atom with a positive charge (mass = 1.673 x 10-24 grams)

How Much Do Atoms Weigh?

Simplest model of an atom

The ability to weigh atoms came about by an observation from an Italian chemist named Amadeo Avogadro. Avogadro was working with gases (nitrogen, hydrogen, oxygen, chlorine) and noticed that when temperature and pressure was the same, these gases combined in definite volume ratios. For example:

One liter of nitrogen combined with three liters of hydrogen to form ammonia (NH3)

One liter of hydrogen combined with one liter of chlorine to make hydrogen chloride (HCl)

Avogadro said that at the same temperature and pressure, equal volumes of the gases had the same number of molecules. So, by weighing the volumes of gases, he could determine the ratios of atomic masses. For example, a liter of oxygen weighed 16 times more than a liter of hydrogen, so an atom of oxygen must be 16 times the mass of an atom of hydrogen. Work of this type resulted in a relative mass scale for elements in which all of the elements related to carbon (chosen as the standard -12). Once the relative mass scale was made, later experiments were able to relate the mass in grams of a substance to the number of atoms and an atomic mass unit (amu) was found; 1 amu or Dalton is equal to 1.66 x 10-24 grams.

At this time, chemists knew the atomic masses of elements and their chemical properties, and an astonishing phenomenon jumped out at them!

The Properties of Elements Showed a Repeating Pattern

At the time that atomic masses had been discovered, a Russian chemist named Dimitri Mendeleev was writing a textbook. For his book, he began to organize elements in terms of their properties by placing the elements and their newly discovered atomic masses in cards. He arranged the elements by increasing atomic mass and noticed that elements with similar properties appeared at regular intervals or periods. Mendeleev's table had two problems:

There were some gaps in his "periodic table."

When grouped by properties, most elements had increasing atomic masses, but some were out of order.

To explain the gaps, Mendeleev said that the gaps were due to undiscovered elements. In fact, his table successfully predicted the existence of gallium and germanium, which were discovered later. However, Mendeleev was never able to explain why some of the elements were out of order or why the elements should show this periodic behavior. This would have to wait until we knew about the structure of the atom.

In the next section, we will look at how we discovered the inside of the atom!

The Structure of the Atom: Early 20th-Century Science

To know the structure of the atom, we must know the following:

What are the parts of the atom?

How are these parts arranged?

Near the end of the 19th century, the atom was thought to be nothing more than a tiny indivisible sphere (Dalton's view). However, a series of discoveries in the fields of chemistry, electricity and magnetism, radioactivity, and quantum mechanics in the late 19th and early 20th centuries changed all of that. Here is what these fields contributed:

How the atom is arranged - quantum mechanics puts it all together: atomic spectra ---> Bohr model of the atom wave-particle duality ---> Quantum model of the atom

Chemistry and Electromagnetism: Discovering the Electron

In the late 19th century, chemists and physicists were studying the relationship between electricity and matter. They were placing high voltage electric currents through glass tubes filled with low-pressure gas (mercury, neon, xenon) much like neon lights. Electric current was carried from one electrode (cathode) through the gas to the other electrode (anode) by a beam called cathode rays. In 1897, a British physicist, J. J. Thomson did a series of experiments with the following results:

By applying an electric field alone, a magnetic field alone, or both in combination, Thomson could measure the ratio of the electric charge to the mass of the cathode rays.

He found the same charge to mass ratio of cathode rays was seen regardless of what material was inside the tube or what the cathode was made of.

Thomson concluded the following:

Cathode rays were made of tiny, negatively charged particles, which he called electrons.

The electrons had to come from inside the atoms of the gas or metal electrode.

Because the charge to mass ratio was the same for any substance, the electrons were a basic part of all atoms.

Because the charge to mass ratio of the electron was very high, the electron must be very small.

Later, an American Physicist named Robert Milikan measured the electrical charge of an electron. With these two numbers (charge, charge to mass ratio), physicists calculated the mass of the electron as 9.10 x 10-28 grams. For comparison, a U.S. penny has a mass of 2.5 grams; so, 2.7 x 1027 or 2.7 billion billion billion electrons would weigh as much as a penny!

Two other conclusions came from the discovery of the electron:

Because the electron was negatively charged and atoms are electrically neutral, there must be a positive charge somewhere in the atom.

Because electrons are so much smaller than atoms, there must be other, more massive particles in the atom.

From these results, Thomson proposed a model of the atom that was like a watermelon. The red part was the positive charge and the seeds were the electrons.

Radioactivity: Discovering the Nucleus, the Proton and the Neutron

Rutherford's view of the atom

About the same time as Thomson's experiments with cathode rays, physicists such as by Henri Becquerel, Marie Curie, Pierre Curie, and Ernest Rutherford were studying radioactivity. Radioactivity was characterized by three types of emitted rays (see How Radioactivity Works for details):

Alpha particles - positively charged and massive. Ernest Rutherford showed that these particles were the nucleus of a helium atom.

The experiment from radioactivity that contributed most to our knowledge of the structure of the atom was done by Rutherford and his colleagues. Rutherford bombarded a thin foil of gold with a beam of alpha particles and looked at the beams on a fluorescent screen, he noticed the following:

Most of the particles went straight through the foil and struck the screen.

Some (0.1 percent) were deflected or scattered in front (at various angles) of the foil, while others were scattered behind the foil.

Rutherford concluded that the gold atoms were mostly empty space, which allowed most of the alpha particles through. However, some small region of the atom must have been dense enough to deflect or scatter the alpha particle. He called this dense region the nucleus (see The Rutherford Experiment for an excellent Java simulation of this important experiment!); the nucleus comprised most of the mass of the atom. Later, when Rutherford bombarded nitrogen with alpha particles, a positively charged particle that was lighter than the alpha particle was emitted. He called these particles protons and realized that they were a fundamental particle in the nucleus. Protons have a mass of 1.673 x 10-24 grams, about 1,835 times larger than an electron!

However, protons could not be the only particle in the nucleus because the number of protons in any given element (determined by the electrical charge) was less than the weight of the nucleus. Therefore, a third, neutrally charged particle must exist! It was James Chadwick, a British physicist and co-worker of Rutherford, who discovered the third subatomic particle, the neutron. Chadwick bombarded beryllium foil with alpha particles and noticed a neutral radiation coming out. This neutral radiation could in turn knock protons out of the nuclei of other substances. Chadwick concluded that this radiation was a stream of neutrally charged particles with about the same mass as a proton; the neutron has a mass of 1.675 x 10-24 grams.

Now that the parts of the atom were known, how were they arranged to make an atom? Rutherford's gold foil experiment indicated that the nucleus was in the center of the atom and that the atom was mostly empty space. So, he envisioned the atom as the positively charged nucleus in the center with the negatively charged electrons circling around it much like a planet with moons. Although he had no evidence that the electrons circled the nucleus, his model seemed reasonable; however, it presented a problem. As the electrons moved in a circle, they would lose energy and give off light. The loss of energy would slow the electrons down. Like any satellite, the slowing electrons would fall into the nucleus. In fact, it was calculated that a Rutherford atom would last only billionths of a second before collapsing! Something was missing!

Quantum Mechanics: Putting It All Together

White light passing through a prism.

Photo courtesy NASA

At the same time that discoveries were being made with radioactivity, physicists and chemists were studying how light interacted with matter. These studies began the field of quantum mechanics and helped solve the structure of the atom.

Quantum Mechanics Sheds Light on the Atom: The Bohr Model

Physicists and chemists studied the nature of the light that was given off when electric currents were passed through tubes containing gaseous elements (hydrogen, helium, neon) and when elements were heated (e.g., sodium, potassium, calcium, etc.) in a flame. They passed the light from these sources through a spectrometer (a device containing a narrow slit and a glass prism).

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Continuous spectrum of white light.

Photo courtesy NASA

Now, when you pass sunlight through a prism, you get a continuous spectrum of colors like a rainbow. However, when light from these various sources was passed through a prism, they found a dark background with discrete lines.

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Hydrogen spectrum

Photo courtesy NASA

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Helium spectrum

Photo courtesy NASA

Each element had a unique spectrum and the wavelength of each line within a spectrum had a specific energy (see How Light Works for details on the relationship between wavelength and energy).

In 1913, a Danish physicist named Niels Bohr put Rutherford's findings together with the observed spectra to come up with a new model of the atom in a real leap of intuition. Bohr suggested that the electrons orbiting an atom could only exist at certain energy levels (i.e., distances) from the nucleus, not at continuous levels as might be expected from Rutherford's model. When atoms in the gas tubes absorbed the energy from the electric current, the electrons became excited and jumped from low energy levels (close to the nucleus) to high energy levels (farther out from the nucleus). The excited electrons would fall back to their original levels and emit energy as light. Because there were specific differences between the energy levels, only specific wavelengths of light were seen in the spectrum (i.e., lines).

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Bohr models of various atoms.

The major advantage of the Bohr model was that it worked. It explained several things:

Atomic spectra - discussed above

Periodic behavior of elements - elements with similar properties had similar atomic spectra.

Each electron orbit of the same size or energy (shell) could only hold so many electrons. For example, the first shell could hold two electrons, the second could hold eight electrons, the third could hold 18 electrons, the fourth 32 and so on until reaching the seventh.

When one shell was filled, electrons were found at higher levels.

Chemical properties were based on the number of electrons in the outermost shell. Elements with full outer shells do not react. Other elements take or give up electrons to get a full outer shell.

As it turns out, Bohr's model is also useful for explaining the behavior of lasers, although these devices were not invented until the middle of the 20th century.

Bohr's model was the predominant model until new discoveries in quantum mechanics were made.

Quantum Mechanics

Branch of physics that deals with the motion of particles by their wave properties at the atomic and subatomic level.

Electrons Can Behave as Waves: The Quantum Model of the Atom

Although the Bohr model adequately explained how atomic spectra worked, there were several problems that bothered physicists and chemists:

Why should electrons be confined to only specified energy levels?

Why don't electrons give off light all of the time? As electrons change direction in their circular orbits (i.e., accelerate), they should give off light.

The Bohr model could explain the spectra of atoms with one electron in the outer shell very well, but was not very good for those with more than one electron in the outer shell.

Why could only two electrons fit in the first shell and why eight electrons in each shell after that? What was so special about two and eight?

Obviously, the Bohr model was missing something!

In 1924, a French physicist named Louis de Broglie suggested that, like light, electrons could act as both particles and waves (see De Broglie Phase Wave Animation for details). De Broglie's hypothesis was soon confirmed in experiments that showed electron beams could be diffracted or bent as they passed through a slit much like light could. So, the waves produced by an electron confined in its orbit about the nucleus sets up a standing wave of specific wavelength, energy and frequency (i.e., Bohr's energy levels) much like a guitar string sets up a standing wave when plucked.

Another question quickly followed de Broglie's idea. If an electron traveled as a wave, could you locate the precise position of the electron within the wave? A German physicist, Werner Heisenberg, answered no in what he called the uncertainty principle:

To view an electron in its orbit, you must shine a wavelength of light on it that is smaller than the electron's wavelength.

We can never know both the momentum and position of an electron in an atom. Therefore, Heisenberg said that we shouldn't view electrons as moving in well-defined orbits about the nucleus!

With de Broglie's hypothesis and Heisenberg's uncertainty principle in mind, an Austrian physicist named Erwin Schrodinger derived a set of equations or wave functions in 1926 for electrons. According to Schrodinger, electrons confined in their orbits would set up standing waves and you could describe only the probability of where an electron could be. The distributions of these probabilities formed regions of space about the nucleus were called orbitals. Orbitals could be described as electron density clouds (see Atomic & Molecular Orbitals for a look at various orbitals). The densest area of the cloud is where you have the greatest probability of finding the electron and the least dense area is where you have the lowest probability of finding the electron.

Wave Functions

Quantum model of a sodium atom.

The wave function of each electron can be described as a set of three quantum numbers:

Principal number (n) - describes the energy level.

Azimuthal number (l) - how fast the electron moves in its orbit (angular momentum); like how fast a CD spins (rpm). This is related to the shape of the orbital.

Magnetic (m) - its orientation in space.

It was later suggested that no two electrons could be in the exact same state, so a fourth quantum number was added. This number was related to the direction that the electron spins while it is moving in its orbit (i.e., clockwise, counterclockwise). Only two electrons could share the same orbital, one spinning clockwise and the other spinning counterclockwise.

The orbitals had different shapes and maximum numbers at any level:

s (sharp) - spherical (max = 1)

p (principal) - dumb-bell shaped (max = 3)

d (diffuse) - four-lobe-shaped (max = 5)

f (fundamental) - six-lobe shaped (max = 7)

The names of the orbitals came from names of atomic spectral features before quantum mechanics was formally invented. Each orbital can hold only two electrons. Also, the orbitals have a specific order of filling, generally:

However, there is some overlap (any chemistry textbook has the details).

The resulting model of the atom is called the quantum model of the atom.

Sodium has 11 electrons distributed in the following energy levels:

one s orbital - two electrons

one s orbital - two electrons and three p orbitals (two electrons each)

one s orbital - one electron

Right now, the quantum model is the most realistic vision of the overall structure of the atom. It explains much of what we know about chemistry and physics. Here are some examples:

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The modern periodic table of the elements (elements are ordered based on atomic number rather than mass).

Chemistry: The Periodic Table - the Table's pattern and arrangement reflects the arrangement of electrons in the atom. Elements have different atomic numbers - the number of protons or electrons increases up the table as electrons fill the shells. Elements have different atomic masses - the number of protons plus neutrons increases up the table. Rows - elements of each row have the same number of energy levels (shells). Columns - elements have the same number of electrons in the outermost energy level or shell (one to eight). Chemical reactions - exchange of electrons between various atoms (giving, taking, or sharing). Exchange involves electrons in the outermost energy level in attempts to fill the outermost shell (i.e., most stable form of the atom).

Electronic components to supply current to the tip, control the scanner and accept the signals from the motion sensor

Computer to control the system and do data analysis (data collection, processing, display)

The STM works like this:

A current is supplied to the tip (probe) while the scanner rapidly moves the tip across the surface of a conducting sample.

When the tip encounters an atom, the flow of electrons between the atom and the tip changes.

The computer registers the change in current with the x,y-position of the atom.

The scanner continues to position the tip over each x,y-point on the sample surface, registering a current for each point.

The computer collects the data and plots a map of current over the surface that corresponds to a map of the atomic positions.

The process is much like an old phonograph where the needle is the tip and the grooves in the vinyl record are the atoms. The STM tip moves over the atomic contour of the surface, using tunneling current as a sensitive detector of atomic position.

The STM and new variations of this microscope allow us to see atoms. In addition, the STM can be used to manipulate atoms as shown here:

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Atoms can be positioned on a surface using the STM tip, creating a custom pattern on the surface.

In summary, science in the 20th century has revealed the structure of the atom. Scientists are now conducting experiments to reveal details of the structure of the nucleus and the forces that hold it together.